10.6 Group 2 Elements : Alkaline Earth Metals

The group 2 elements comprise beryllium, magnesium, calcium, strontium, barium and radium. They follow alkali metals in the periodic table. These (except beryllium) are known as alkaline earth metals.

The first element beryllium differs from the rest of the members and shows diagonal relationship to aluminium.

 

10.6.1 Electronic Configuration

These elements have two electrons in the s -orbital of the valence shell (Table 10.2). Their general electronic configuration may be represented as [noble gas] ns2.

Like alkali metals, the compounds of these elements are also predominantly ionic.

 

10.6.2 Atomic and Ionic Radii

The atomic and ionic radii of the alkaline earth metals are smaller than those of the corresponding alkali metals in the same periods. This is due to the increased nuclear charge in these elements.

Within the group, the atomic and ionic radii increase with increase in atomic number.

 

10.6.3 Ionization Enthalpies

The alkaline earth metals have low ionization enthalpies due to fairly large size of the atoms. Since the atomic size increases down the group, their ionization enthalpy decreases.

The first ionisation enthalpies of the alkaline earth metals are higher than those of the corresponding Group 1 metals. This is due to their small size as compared to the corresponding alkali metals.

It is interesting to note that the second ionisation enthalpies of the alkaline earth metals are smaller than those of the corresponding alkali metals.

 

10.6.4 Hydration Enthalpies

Like alkali metal ions, the hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group.

Be2+> Mg2+ > Ca2+ > Sr2+ > Ba2+

The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions.

Thus, compounds of alkaline earth metals are more extensively hydrated than those of alkali metals, e.g., MgCl2 and CaCl2 exist as MgCl2.6H2O and CaCl2· 6H2O while NaCl and KCl do not form such hydrates.

 

10.6.5 Physical Properties

The alkaline earth metals, in general, are silvery white, lustrous and relatively soft but harder than the alkali metals. Beryllium and magnesium appear to be somewhat greyish.

The melting and boiling points of these metals are higher than the corresponding alkali metals due to smaller sizes. The trend is, however, not systematic.

Because of the low ionisation enthalpies, they are strongly electropositive in nature. The electropositive character increases down the group from Be to Ba. Calcium, strontium and barium impart characteristic brick red, crimson and apple green colours respectively to the flame.

In flame the electrons are excited to higher energy levels and when they drop back to the ground state, energy is emitted in the form of visible light.

The electrons in beryllium and magnesium are too strongly bound to get excited by flame. Hence, these elements do not impart any colour to the flame.

The flame test for Ca, Sr and Ba is helpful in their detection in qualitative analysis and estimation by flame photometry. The alkaline earth metals like those of alkali metals have high electrical and thermal conductivities which are typical characteristics of metals.

 

10.6.6 Chemical Properties

The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group.

Reactivity towards air and water: Beryllium and magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface.

However, powdered beryllium burns brilliantly on ignition in air to give BeO and Be3N2. Magnesium is more electropositive and burns with dazzling brilliance in air to give MgO and Mg3N2.

Calcium, strontium and barium are readily attacked by air to form the oxide and nitride. They also react with water with increasing vigour even in cold to form hydroxides.

Reactivity towards the halogens: All the alkaline earth metals combine with halogen at elevated temperatures forming their halides.

M +X2 → MX2 (X = F, Cl, Br, I)

Thermal decomposition of (NH4)2BeF4 is the best route for the preparation of BeF2, and BeCl2 is conveniently made from the oxide.

BeO + C +Cl2 → BeCl2 + CO

Reactivity towards hydrogen: All the elements except beryllium combine with hydrogen upon heating to form their hydrides, MH2.

BeH2, however, can be prepared by the reaction of BeCl2 with LiAlH4.

2BeCl2 + LiAlH4 → 2BeH2 + LiCl + AlCl3

Reactivity towards acids: The alkaline earth metals readily react with acids liberating dihydrogen.

M + 2HCl → MCl2 + H2

Reducing nature: Like alkali metals, the alkaline earth metals are strong reducing agents. This is indicated by large negative values of their reduction potentials.

However their reducing power is less than those of their corresponding alkali metals. Beryllium has less negative value compared to other alkaline earth metals.

However, its reducing nature is due to large hydration energy associated with the small size of Be2+ ion and relatively large value of the atomization enthalpy of the metal.

Solutions in liquid ammonia: Like alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.

M +(x+y) NH3 → [M(NH3)x]2+ + 2[e(NH3)y]

From these solutions, the ammoniates, [M(NH3)6]2+ can be recovered.

 

10.6.7 Uses

Beryllium is used in the manufacture of alloys. Copper-beryllium alloys are used in the preparation of high strength springs. Metallic beryllium is used for making windows of X-ray tubes.

Magnesium forms alloys with aluminium, zinc, manganese and tin. Magnesium-aluminium alloys being light in mass are used in air-craft construction. Magnesium (powder and ribbon) is used in flash powders and bulbs, incendiary bombs and signals.

A suspension of magnesium hydroxide in water (called milk of magnesia) is used as antacid in medicine. Magnesium carbonate is an ingredient of toothpaste.

Calcium is used in the extraction of metals from oxides which are difficult to reduce with carbon. Calcium and barium metals, owing to their reactivity with oxygen and nitrogen at elevated temperatures, have often been used to remove air from vacuum tubes.

Radium salts are used in radiotherapy, for example, in the treatment of cancer.

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